Difficulty: Introductory to Intermediate | Prerequisites: Know the difference between molecular and ionic compounds (see Matter, Substances, and Classification notes).
Naming compounds correctly is the language of chemistry. If you cannot convert between a formula and a name, you will struggle with every reaction, equation, and lab report from here on. This section also introduces three ways to write a formula (molecular, empirical, structural) and the mole, which is the bridge between the atomic scale and the lab scale. You should already know the difference between molecular and ionic compounds before starting here.
Molecular compounds are named using Greek prefixes (mono-, di-, tri-, etc.) with the second element ending in "-ide." Ionic compounds use no prefixes: name the metal, then the non-metal with an "-ide" ending. A molecular formula shows the exact number of each atom; an empirical formula shows the simplest whole-number ratio; a structural formula shows how atoms are connected. One mole = 6.022 x 10²³ particles (Avogadro's number), and it links atomic mass units to grams.
Molecular formula
Shows the exact number of atoms of each element in one molecule. Example: acetic acid = C₂H₄O₂.
Empirical formula
The simplest whole-number ratio of atoms in a compound. Example: acetic acid's empirical formula is CH₂O (dividing C₂H₄O₂ by 2).
Think of it as: the molecular formula reduced to its lowest terms, the way you simplify a fraction.
Structural formula
Shows how the atoms are connected to each other with lines representing bonds. It reveals the arrangement, not just the count.
Atomic mass unit (amu)
Defined as 1/12 the mass of a carbon-12 atom. 1 amu = 1.993 x 10⁻²³ g. This is the scale used for individual atoms and molecules.
Mole (mol)
The SI unit for "amount of substance." One mole contains exactly 6.022 x 10²³ particles (atoms, molecules, ions, or whatever you are counting). This number is called Avogadro's number.
In simple terms, a mole is just a very large counting number, like a "dozen" is 12. One mole of carbon-12 atoms weighs exactly 12 grams.
Avogadro's number (Nₐ)
6.022 x 10²³. The number of particles in one mole of any substance.
Nomenclature prefixes (for molecular compounds)
Greek prefixes indicating the number of atoms: mono- (1), di- (2), tri- (3), tetra- (4), penta- (5), hexa- (6), hepta- (7), octa- (8), nona- (9), deca- (10).
-ide suffix
The ending applied to the second element in a binary compound's name. Oxygen becomes "oxide," chlorine becomes "chloride," and so on.
The rule: name the first element, then the second element with the suffix "-ide." Use Greek prefixes to indicate how many atoms of each element are present.
CO₂ = carbon dioxide
HCl = hydrogen chloride
SCl₄ = sulphur tetrachloride (note: "tetrafluoride" in some texts)
SiC = silicon carbide
OF₂ = oxygen difluoride
PCl₅ = phosphorus pentachloride
N₂O₄ = dinitrogen tetroxide
NI₃ = nitrogen triiodide
H₂S = hydrogen sulphide
The prefix "mono-" is usually dropped for the first element. You would not say "monocarbon dioxide"; you say "carbon dioxide."
Some molecular compounds have traditional names that override the prefix system. You must memorise these:
H₂O = water
NH₃ = ammonia
CH₄ = methane
PH₃ = phosphine
The rule: name the metal first, then the non-metal with the suffix "-ide." Never use Greek prefixes for ionic compounds. The ratio is determined by the charges of the ions, not by prefixes.
MgCl₂ = magnesium chloride (Mg²⁺ needs two Cl⁻ to balance)
NaCl = sodium chloride
Molecular formula: the actual atom count. Acetic acid = C₂H₄O₂.
Empirical formula: the simplest ratio. Acetic acid = CH₂O (divide all subscripts by 2).
Structural formula: a drawing showing how atoms are bonded. It reveals the arrangement (e.g. that acetic acid has a -COOH group).
Subscripts in a formula also tell you the mole ratio. In H₂O, there are 2 moles of hydrogen for every 1 mole of oxygen.
1 amu = 1/12 the mass of a carbon-12 atom = 1.993 x 10⁻²³ g.
1 mole = 6.022 x 10²³ particles (Avogadro's number).
The mole bridges the atomic world and the lab world: the atomic mass of an element in amu, when taken in grams, gives you exactly one mole of that element. Carbon has an atomic mass of 12.01 amu, so 12.01 g of carbon = 1 mole of carbon atoms.
Subscripts in a chemical formula tell you the mole ratio between elements in that compound.
Nomenclature is not just academic. Drug names, material science labels, and environmental regulations all rely on systematic chemical names. If a lab report says "dinitrogen tetroxide," anyone in chemistry knows the exact formula (N₂O₄) without ambiguity.
The mole is the unit that makes chemistry practical. You cannot weigh out individual atoms, but you can weigh out 12.01 grams of carbon and know you have 6.022 x 10²³ atoms of it.
Students often use prefixes when naming ionic compounds. Never do this. MgCl₂ is "magnesium chloride," not "magnesium dichloride." The ratio is implied by the charges.
Students confuse molecular and empirical formulas. C₂H₄O₂ and CH₂O are not the same compound in general; the empirical formula just gives the ratio. Different compounds can share the same empirical formula.
Students sometimes think a mole is a mass. It is not. A mole is a count (6.022 x 10²³ items). The mass of one mole differs for every substance.
Students forget the common-name exceptions (water, ammonia, methane, phosphine) and try to name them with prefixes. "Dihydrogen monoxide" is a joke, not a real chemical name.
⚠️ Know the Greek prefixes from mono- through deca-. These will appear in naming questions.
⚠️ Be able to name both molecular and ionic compounds from their formula and vice versa.
⚠️ Memorise the four common-name exceptions: H₂O (water), NH₃ (ammonia), CH₄ (methane), PH₃ (phosphine).
⚠️ Be ready to convert between molecular and empirical formulas.
⚠️ Understand that 1 mole = 6.022 x 10²³ and that the molar mass in grams numerically equals the atomic/molecular mass in amu.
⚠️ The lecture references Exam 1 2009 questions 15-18 and OWL 4.4d as practice resources.
Fill in the blank: ionic compounds are named without ______.
True or false: the empirical formula of C₆H₁₂O₆ is CH₂O.
Fill in the blank: one mole contains ______ particles.
Name the compound N₂O₄.
True or false: NH₃ is named "nitrogen trihydride" in standard nomenclature.
Answers: 1. Prefixes. 2. True (6:12:6 simplifies to 1:2:1). 3. 6.022 x 10²³. 4. Dinitrogen tetroxide. 5. False (it is called ammonia, by convention).
Q: Name the compound PCl₅ and classify it as molecular or ionic.
A: Phosphorus pentachloride. Molecular (both phosphorus and chlorine are non-metals).
Q: Write the formula for magnesium chloride and explain why no prefix is used.
A: MgCl₂. It is an ionic compound (Mg is a metal, Cl is a non-metal), and ionic compounds are named using the charges of the ions, not Greek prefixes. Mg²⁺ requires two Cl⁻ to balance.
Q: The molecular formula of a compound is C₄H₈O₂. What is its empirical formula?
A: C₂H₄O. Divide all subscripts by 2.
Q: How many molecules are in 2 moles of water?
A: 2 x 6.022 x 10²³ = 1.204 x 10²⁴ molecules.
Q: A compound's empirical formula is CH₂O and its molar mass is approximately 180 g/mol. What is its molecular formula?
A: The empirical formula mass of CH₂O is 12 + 2(1) + 16 = 30 g/mol. 180 / 30 = 6. So the molecular formula is C₆H₁₂O₆ (glucose).
Nomenclature is the starting point for writing balanced chemical equations: you need to know what a name means to write its formula. The mole concept connects directly to stoichiometry, where you convert between grams, moles, and molecules to predict how much product a reaction will yield. Empirical and molecular formulas reappear in combustion analysis problems.
Chemical nomenclature, naming compounds, molecular compound naming, ionic compound naming, Greek prefixes, -ide suffix, molecular formula, empirical formula, structural formula, mole, Avogadro's number, amu, atomic mass unit, molar mass, CHEM 101, general chemistry, Michigan State University